Compare the Lewis structures of CF₄ and NF₃ Why do these molecule...

Question

Compare the Lewis structures of CF₄ and NF₃ Why do these molecules have different shapes?

Accepted Answer

Welcome back, everyone. Based on their le structures explain why carbon tetrachloride and nitrogen trichloride have different molecular shapes. Choice. A says that because carbon tetrachloride has a tetrahedral shape since the central atom, carbon has four bonded atoms being chlorine. While nitrogen trichloride has a trigonal pure middle shape since its central atom, nitrogen has three bonded atoms being chlorine and one lone electron pair. Choice B says that because carbon tetrachloride has a tetrahedral shape since the central atom, carbon has four bonded atoms being chlorine. While nitrogen trichloride has a trigonal planar shape. Since the central atom, nitrogen has three bonded atoms being chlorine. Choice C says carbon tetrachloride has a trigonal pyramidal shape. Since the central atom, carbon has three bonded atoms being chlorine and one lone electron pair. While nitrogen trichloride has a tetrahedral shape, since the central atom, nitrogen has four bonded atoms being chlorine. And choice D states that carbon tetrachloride has a trigonal planar shape. Since the central atom carbon has three bonded atoms being chlorine. While nitrogen trichloride has a tetrahedral shape. Since the central atom, nitrogen has four bonded atoms being chlorine. So we're going to need to determine which which of our answer choices is correct by drawing out the Lewis structure of carbon tetrachloride as well as nitrogen trichloride. But let's begin with carbon tetrachloride. We're going to calculate total valence electrons for carbon tetrachloride. Beginning with carbon being a group four A element will correspond to four valence electrons. Next, for chlorine, we have chlorine as a group five A element on our product table Correction A Group seven A element on our periodic table corresponding to I'm sorry. This says seven a. So group seven a corresponds to seven valence electrons. And because we have a subscript of four in our carbon tetrachloride chemical formula, we would multiply by a total of four chlorine atoms which will yield a contribution of 28 valence electrons via chlorine. So adding up our total, we have 28 plus four which gives us a total of 32 valence electrons total for our structure. So now placing our single carbon in the center and surrounding it by our four chlorine atoms, we're going to begin by recalling the bonding preference of carbon being in group four A with its four valence electrons, it's going to prefer to form four bonds, two other atoms. So we would form four bonds to our four chlorine atoms via our central l carbon atoms. Sorry. And this will use up a total of eight of our 32 vals electrons leaving us with a balance of 24 valence electrons remaining. And Next, we're going to consider our chlorine, which currently only have one directly attached val electron, which is being covalent shared in the bond to carbon. So we're going to need to complete the advanced electrons on chlorine by adding a set of three lone pairs to each of our chlorine atoms. So drawing those in, we would have again, three lone pairs added to each of our four chlorine atoms in our structure, Which will use up a total of our 24 remaining balance electrons leaving us with a balance of zero. And we would confirm that we have a net formal charge of zero for the structure. So we have the best Luis structure for carbon tetrachloride. We're going to account for electron groups bonding groups and lone pairs to our central atom. So beginning with bonding groups, we observe four chlorine atoms bonded to our central atom carbon, accounting for four bonding groups total. And we would observe that we have zero lone pairs on our central atom carbon. So for our total number of electron groups, we would have four because we have four electron groups, four being bonding and zero being lone pairs, we would therefore have a molecular geometry which is going to be tetrahedral. And we have a perfect tetrahedral shape because we have a symmetrical molecule where we are bonded to four chlorine atoms and have no lone pairs influencing the shape of how these atoms are bonded. And so thus, we have a symmetrical tetrahedral perfect shape. Now, we're going to compare to our les structure of nitrogen trichloride and seal sub three calculating total valence electrons. First, beginning with nitrogen, which we recognize on our periodic table is located in group five A corresponding to five valence electrons moving on to chlorine, which we recall from earlier is located in group seven A on our periodic table. Our halogen group corresponding to seven valence electrons. And given by the subscript of three, we multiply by three atoms of chlorine which will yield a contribution of 21 valence electrons via chlorine. So adding the total, we have 21 plus five, which is going to give us a total 26 valence electrons in our structure. So placing our single nitrogen in the center and surrounding it by our three chlorine atoms, we're going to begin by recalling our bonding preference of nitrogen, which is to be bonded to three atoms. So we'll form three covalent bonds to chlorine and having a lone pair on itself. So drawing in that lone pair at the top of nitrogen, we would have used up a total of eight of our 26 bounce electrons leaving us with a balance of 18 vals electrons remaining in which now we're going to consider our only single vals electron directly attached to each chlorine in our structure, which is shared in the covalent bonds to our central atom nitrogen. And recall that chlorine requires seven valence electrons to have a formal charge of zero and to be stable. So we're going to need to fill in a total of three sets of lone pairs on each of our chlorine atoms. And this is going to use up a total of our 18 remaining valence electrons leaving us with a balance of zero valence electrons giving us a net formal charge of zero for this structure. Meaning we have the best lowest structure for nitrogen trichloride. So now considering electron groups bonding groups and then lone pairs on our central atom nitrogen, beginning with bonding groups, we would observe our three chlorine atoms bonded to our central atom nitrogen. So that's three bonding groups. We have one lone pair on our nitrogen at the top counting for one lone pair on our central atom nitrogen. So our three bonding groups and the one loan pair account for a total of four electron groups. And based on the presence of this lone pair and our three bonded groups, we have an influence on our shape of our molecule, which is now going to result in a molecular geometry which is trigonal pure middle. Again, due to the presence of this lone pair, affecting how our chlorine will now be bonded. So the shape of our molecule is more accurately drawn in a shape like this where our lone pair is at the top of our central atom. So now going back to our answer choices. Now that we understand the molecular shapes of carbon tetrachloride and nitrogen trichloride, we would rule out choice b since it does not state the fact that we have that lone electron pair influencing the three, the way that our three nitrogen atoms are bonded in nitrogen trichloride. We would also rule out choice C which states that carbon tetrachloride has a lone pair, which is incorrect. It does not have a lone pair in its structure. Next, we have choice d that we would rule out because it incorrectly states that nitrogen trichloride has a tetrahedral shape. So this leaves us with choice. A which states that carbon tetrachloride has a tetrahedral shape since the central atom, carbon has four bonded atoms being chlorine, while nitrogen trichloride has a trigonal pyramidal shape. Since the central atom, nitrogen has three bonded atoms being chlorine and one lone electron pair. This is our final answer. A as to why carbon tetrachloride and nitrogen trichloride have different molecular shapes. I hope this made sense. And let us know if you have any questions.

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Compare the Lewis structures of CF₄ and NF₃ Why do these molecules have different shapes?

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